Chapter 1Carbon Compounds and Chemical

Bonds

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What is organic chemistry?In general, Organic chemistry is the study of the compounds of carbon.

The compounds of carbon are the central substances of which: all living things include DNA, proteins and enzymes.

in addition, We live in an Age of organic chemistry: The clothing we wear, whether a natural substance (wool or cotton) or a

synthetic (nylon or polyester), is made up of carbon compound Most of the medicines that help us cure diseases and relieve suffering

are organic Most electronic and house equipment for example TV, mobile,

furniture, .. contains rubber or plastic, both are organic compound

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Structural TheoryTow central ideas make up The structural theory

1. Valency: the atoms of the elements in organic compounds can form a fixed number of bonds

2. A Carbon atom can form one or more bonds to other carbons

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Isomerism Isomers

are different compounds that have the same molecular formula. There is many types of isomeric compounds are exist in life.

For Example:

The Constitutional Isomers • Constitutional isomers are one type of isomer.• They are different compounds that have the same molecular

formula but different connectivity of atoms (in the way in which their atoms are bonded together)

• They often differ in physical properties (e.g. boiling point, melting point, density) and chemical properties

Consider two compounds with molecular formula C2H6O

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H C C

H

H H

H

OH

Ethayl Alcohol

Boiling point = 78.5 oC

liquid

H C O

H

H

C

H

H

H

Dimethyl Ether

Boiling point = -24.9 oC

Gas

Constitutional isomersmolecular forula C2H6O

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Three Dimensional Shape of Molecules Virtually all molecules possess a 3-dimensional shape which is

often not accurately represented by drawings.

It was proposed in 1874 by van’t Hoff and le Bel that the four bonds around carbon where not all in a plane but rather in a tetrahedral arrangement i.e. the four C-H bonds point towards the corners of a regular tetrahedron

Bond in plane of paperBond going back into paper away from you

Bond coming out of paper towards you

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Chemical Bonds

The Octet RuleOctet Rule

Atoms form bonds to produce the electron configuration of a noble gas (because the electronic configuration of noble gases is particularly stable)

For most atoms of interest this means achieving a valence shell configuration of 8 electrons corresponding to that of the nearest noble gas

Atoms close to helium achieve a valence shell configuration of 2 electrons

To satisfy the octet rule

Atoms can transfer of Atoms can share electros to

one or more electrons create covalent bonds

from one atom to another

to create ions ionic bond

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Electronegativity

Electronegativity is the ability of an atom to attract electrons

It increases from left to right and from bottom to top in the periodic table (noble gases excluded)

Fluorine is the most electronegative atom and can stabilize excess electron density the best

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Ionic BondsWhen ionic bonds are formed atoms gain or lose electrons to achieve the

electronic configuration of the nearest noble gas

In the process the atoms become ionic The resulting oppositely charged ions attract and form ionic bonds

This generally happens between atoms of widely different electronegativities

Example Lithium loses an electron (to have the configuration of helium) and

becomes positively charged Fluoride gains an electron (to have the configuration of neon) and

becomes negatively charged The positively charged lithium and the negatively charged fluoride

form a strong ionic bond (actually in a crystalline lattice)

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Covalent Bonds Covalent bonds occur between atoms of similar electronegativity

(close to each other in the periodic table) Atoms achieve octets by sharing of valence electrons Molecules result from this covalent bonding Valence electrons can be indicated by dots (electron-dot formula

or Lewis structures). The usual way to indicate the two electrons in a bond is to use a

line (one line = two electrons)

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Polar bondC-Cl

Non polar bondC-C

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Atoms can be assigned electronic configuration using the following rules: Aufbau Principle: The lowest energy orbitals are filled first Pauli Exclusion Principle: A maximum of two spin paired

electrons may be placed in each orbital Hund’s Rule: One electron is added to each degenerate (equal

energy orbital) before a second electron is added

Electronic Configurations of Some Second Row Elements

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Molecular Orbitals (MOs) A simple model of bonding is illustrated by forming molecular H2

from H atoms and varying distance: Region I: The total energy of two isolated atoms Region II: The nucleus of one atom starts attracting the electrons of the other; the

energy of the system is lowered Region III: at 0.74 Å the attraction of electrons and nuclei exactly balances

repulsion of the two nuclei; this is the bond length of H2

Region IV: energy of system rises as the repulsion of the two nuclei predominates

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Bonding Molecular Orbitals (molec) AOs combine by addition (the AOs of the same phase sign

overlap) The wave functions reinforce The value of increases between the two nuclei The value of 2 (electron probability density) in the region

between the two nuclei increases The two electrons between the nuclei serve to attract the nuclei

towards each other This is the ground state (lowest energy state) of the MO

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The Structure of Methane and Ethane: sp3 Hybridization

The structure of methane with its four identical tetrahedral bonds cannot be adequately explained using the electronic configuration of carbon

Hybridization of the valence orbitals (2s and 2p) provides four new identical orbitals which can be used for the bonding in methane

Orbital hybridization is a mathematical combination of the 2s and 2p wave functions to obtain wave functions for the new orbitals

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When one 2s orbital and three 2p orbitals are hybridized four new and identical sp3 orbitals are obtained

When four orbitals are hybridized, four orbitals must result

Each new orbital has one part s character and 3 parts p character

The four identical orbitals are oriented in a tetrahedral arrangements

The antibonding orbitals are not derived in the following diagram

The four sp3 orbitals are then combined with the 1s orbitals of four hydrogens to give the molecular orbitals of methane

Each new molecular orbital can accommodate 2 electrons

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An sp3 orbital looks like a p orbital with one lobe greatly extended Often the small lobe is not drawn

The extended sp3 lobe can then overlap well with the hydrogen 1s to form a strong bond

The bond formed is called a sigma () bond because it is circularly symmetrical in cross section when view along the bond axis

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A variety of representations of methane show its tetrahedral nature and electron distribution

a. calculated electron density surface b. ball-and-stick model c. a typical 3-dimensional drawing

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Methane CH4

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Ethane (C2H6) The carbon-carbon bond is made from overlap of two sp3 orbitals

to form a bond The molecule is approximately tetrahedral around each carbon

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The representations of ethane show the tetrahedral arrangement around each carbon

a. calculated electron density surface b. ball-and-stick model c. typical 3-dimensional drawing

Generally there is relatively free rotation about bonds Very little energy (13-26 kcal/mol) is required to rotate around the carbon-carbon

bond of ethane

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The Structure of Ethene (Ethylene) : sp2 Hybridization

Ethene (C2H2) contains a carbon-carbon double bond and is in the class of organic compounds called alkenes

Another example of the alkenes is propene

The geometry around each carbon is called trigonal planar All atoms directly connected to each carbon are in a plane The bonds point towards the corners of a regular triangle The bond angle are approximately 120o

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There are three bonds around each carbon of ethene and these are formed by using sp2 hybridized orbitals

The three sp2 hybridized orbitals come from mixing one s and two p orbitals

One p orbital is left unhybridized

The sp2 orbitals are arranged in a trigonal planar arrangement The p orbital is perpendicular (orthoganol) to the plane

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Overlap of sp2 orbitals in ethylene results in formation of a framework

One sp2 orbital on each carbon overlaps to form a carbon-carbon bond; the remaining sp2 orbitals form bonds to hydrogen

The leftover p orbitals on each carbon overlap to form a bonding bond between the two carbons

bond results from overlap of p orbitals above and below the plane of the bond

It has a nodal plane passing through the two bonded nuclei and between the two lobes of the molecular orbital

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Restricted Rotation and the Double Bond There is a large energy barrier to rotation (about 264 kJ/mol)

around the double bond This corresponds to the strength of a bond The rotational barrier of a carbon-carbon single bond is 13-26 kJ/mol

This rotational barrier results because the p orbitals must be well aligned for maximum overlap and formation of the bond

Rotation of the p orbitals 90o totally breaks the bond

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Cis-trans isomers Cis-trans isomers are the result of restricted rotation about double

bonds These isomers have the same connectivity of atoms and differ

only in the arrangement of atoms in space This puts them in the broader class of stereoisomers

The molecules below do not superpose on each other One molecule is designated cis (groups on same side) and the

other is trans (groups on opposite side)

Cis-trans isomerism is not possible if one carbon of the double bond has two identical groups

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The Structure of Ethyne (Acetylene): sp Hybridization

Ethyne (acetylene) is a member of a group of compounds called alkynes which all have carbon-carbon triple bonds

Propyne is another typical alkyne

The arrangement of atoms around each carbon is linear with bond angles 180o

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The carbon in ethyne is sp hybridized One s and one p orbital are mixed to form two sp orbitals Two p orbitals are left unhybridized

The two sp orbitals are oriented 180o relative to each other around the carbon nucleus

The two p orbitals are perpendicular to the axis that passes through the center of the sp orbitals

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In ethyne the sp orbitals on the two carbons overlap to form a bond

The remaining sp orbitals overlap with hydrogen 1s orbitals

The p orbitals on each carbon overlap to form two bonds The triple bond consists of one and two bonds

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Bond Lengths of Ethyne, Ethene and Ethane The carbon-carbon bond length is shorter as more bonds hold the

carbons together With more electron density between the carbons, there is more “glue” to hold the

nuclei of the carbons together

The carbon-hydrogen bond lengths also get shorter with more s character of the bond

2s orbitals are held more closely to the nucleus than 2p orbitals A hybridized orbital with more percent s character is held more closely to the

nucleus than an orbital with less s character The sp orbital of ethyne has 50% s character and its C-H bond is shorter The sp3 orbital of ethane has only 25% s character and its C-H bond is longer

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The Effect of Hybridization on AcidityHydrogens connected to orbitals with more s character

will be more acidic s orbitals are smaller and closer to the nucleus than p orbitals

Anions in hybrid orbitals with more s character will be held more closely to the nucleus and be more stabilized and less acidic

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Summery The number of molecular orbitals formed equals the

number of the atomic orbitals used Hybridized orbitals are obtained by mixing the wave

functions of different types of orbitals Four sp3 orbitals are obtained from mixing one s and three p orbitals

The geometry of the four orbitals is tetrahedral This is the hybridization used in the carbon of methane

Three sp2 orbitals are obtained from mixing one s and two p orbitals The geometry of the three orbitals is trigonal planar The left over p orbital is used to make a bond This is the hybridization used in the carbons of ethene

Two sp orbitals are obtained from mixing one s and one p orbital The geometry of the two orbitals is linear The two leftover p orbitals are used to make two bonds This is the hybridization used in the carbons of ethyne

Sigma ( bonds have circular symmetry when viewed along the bond axis

Pi () bonds result from sideways overlap of two p orbitals

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Representations of Structural Formulas Dot formulas are more cumbersome to draw than dash formulas

and condensed formulas Lone-pair electrons are often (but not always) drawn in, especially

when they are crucial to the chemistry being discussed

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Dash formulas Each dash represents a pair of electrons This type of representation is meant to emphasize connectivity

and does not represent the 3-dimensional nature of the molecule The dash formulas of propyl alcohol appear to have 90o angles for carbons which

actually have tetrahedral bond angles (109.5o)

There is relatively free rotation around single bonds so the dash structures below are all equivalent

Constitutional isomers Constitutional isomers have the same molecular formula but different connectivity Propyl alcohol (above) is a constitutional isomer of isopropyl alcohol (below)

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Condensed Structural Formulas In these representations, some or all of the dash lines are omitted In partially condensed structures all hydrogens attached to an

atom are simply written after it but some or all of the other bonds are explicitly shown

In fully condensed structure all bonds are omitted and atoms attached to carbon are written immediately after it

For emphasis, branching groups are often written using vertical lines to connect them to the main chain

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Bond-Line Formulas A further simplification of drawing organic molecules is to

completely omit all carbons and hydrogens and only show heteroatoms (e.g. O, Cl, N) explicitly

Each intersection or end of line in a zig-zag represents a carbon with the appropriate amount of hydrogens

Heteroatoms with attached hydrogens must be drawn in explicitly

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Cyclic compounds are condensed using a drawing of the corresponding polygon

Multiple bonds are indicated by using the appropriate number of lines connecting the atoms

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Three-Dimensional Formulas Since virtually all organic molecules have a 3-dimensional shape it is

often important to be able to convey their shape

The conventions for this are: Bonds that lie in the plane of the paper are indicated by a simple line Bonds that come forward out of the plane of the paper are indicated by a solid wedge Bonds that go back out of the plane of the paper are indicated by a dashed wedge

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Generally to represent a tetrahedral atom:Two of the bonds are drawn in the plane of the paper about 109o apartThe other two bonds are drawn in the opposite direction to the in- plane bonds but right next to each other

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Trigonal planar arrangements of atoms can be drawn in 3-dimensions in the plane of the paper

Bond angles should be approximately 120o

These can also be drawn side-on with the central bond in the plane of the paper, one bond forward and one bond back

Linear arrangements of atoms are always best drawn in the plane of the paper

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